STRUCTURES OF ALUMINUM HYDROXIDE AND · PDF filethe american mineralogist, vol. 55,...

THE AMERICAN MINERALOGIST, VOL. 55, JANUARY-FEBRUARY, i97O

STRUCTURES OF ALUMINUM HYDROXIDE ANDGEOCHEMICAL IMPLICATIONS1

Rononr ScnoBN aNo Cuenr.Bs E. RonBnsoN, U.S. Geo-logical Suntey, Menlo Park, CaliJornia 94025.

AgsrnacrSynthesis experiments in the alumina-water system at room temperature indicate that

the gibbsite polymorph precipitates slowly from solutions whose pH is below the point ofminimum solubility (pH 5.8), and the bayerite polymorph precipitates rapidly from solu-tions whose pH is above 5.8. Nordstrandite, the third polymorph of aluminum hydroxide,forms from bayerite during aging at intermediate to high pH values. fn solutions of inter-mediate pH, both gibbsite and bayerite form, but with aging, early-formed gibbsite disap-pears as more bayerite forms During aging, the pH's of the mother liquors decrease ifgibbsite precipitates and increase if bayerite precipitates.

The principal structural difierence among the three aluminum hydroxide polyrnorphsis in the mode of stacking successive layers. This, in turn, is controlled by the shape of,and charge distribution on, hydroxyl ions located at the surfaces of the layers. The extentof polarization by the aluminum cations determines whether the hydroxyl ions will possesscy'lindrical symmetry as in bayerite or tetrahedral symmetry as in gibbsite. We suggest thatthe extent of polarization of the hydroxyl ions within the solid phases is inherited from thepolarization attained by the hydroxyl ions in aqueous aluminum complexes in the motherliquors.

The usual abundance of the silicate ion in natural alkaline environments may explainthe scarcity of nordstrandite, for the silicate ion may favor the precipitation of alumino-silicate minerals rather than aluminum hydroxide. Near absence of bayerite in naturemay indicate that it is metastable and will invert to nordstrandite in alkaline solutions.

INrnoluctroN

The role piayed by aluminum during weathering of the earth's crustis an enigma. The claim by Rankama and Sahama (1950) that "Thecycle of aluminum is simple, and its details are well known." is no longeradequate. As our understanding of the geochemistry of other commonelements improves, gaps in our knowledge of the geochemistry of alumi-num stand out.

For example, the presence of abundant authigenic aluminosilicateminerals in sedimentary rocks means that large quantities of the con-stituent elements were in solution. Aluminum is one of the few elementswhose solubility is so low in most natural waters as to appear incapableof taking part in the formation of vast amounts of authigenic alumino-silicates. Similarly, the synthesis of aluminosilicates in the laboratory atroom temperature and under conditions resembling natural environ-ments is not easily attained. Bauxite deposits, the primary ore of alu-minum, form by residual concentration of hydrous aluminum oxidesduring weathering of aluminosilicate minerals. We infer that silica andother cations, with the exception of aluminum, go into solution. But this

I Publication authorized by the Director, U. S. Geological Survey.

43

ROBERT SCHOEN AND CHARLES E, ROBERSON

mechanism fails to explain bauxite deposits formed on exceedingly low-

aluminous rocks. Nor does the common presence of concretionary piso-

Iites in bauxite deposits strengthen our conviction in the absolute im-

mobility of aluminum during weathering. The dogma of aluminum's low

solubility at intermediate pH values leads biologists to propose that

aluminum is nonessential in life processes in spite of its wide distribution

in organisms (Hutchinson, 1945).Our principal gaps in understanding the geochemistry of aluminum

arise from the lack of detailed knowledge of the controls on solubility as

well as the kinds and amounts of substances in solution. In addition, the

aluminous solids that precipitate from supersaturated solutions must be

adequately characterized. A fuller knowledge of the geochemistry of

aluminum will help clarify the origins of the most abundant minerals in

crustal rocks, as well as form an acceptable theory of genesis for bauxite

ore deposits, and perhaps even shed Iight on critical economic problems

surrounding the formation of alumina catalysts.Recent studies of the alumina-water system at room temperature by

the U. S. Geological Survey aim to clarify and fill some of these gaps in

knowledge. This paper, an outgrowth of the studies by Hem and Rober-

son (1967) and Hem (1968), and of their continuing research into the

nature of aluminum in natural waters, describes the problems of identi-

fi.cation of small amounts of fine-grained synthetic aluminum hydroxideprecipitates formed in dilute solutions. The results of our X-ray diffrac-

tion study provide a basis for suggesting a mechanism controlling the

formation of various polymorphs of aluminum hydroxide. An explanation

of the distribution in nature of the minerals gibbsite, nordstrandite, and

bayerite follows as a consequence of the application of this mechanism.

Puon Svnrnnsrs Srurps

There has been extensive research for many years on factors that control the formation

of compounds in the alumina-water system in the laboratory. We give here only a summary

of the most recent work as background for the part of this paper dealing with synthesis of

aluminum hydroxide.

A chapter by Rooksby (1961) on the oxides and hydroxides of aluminum and iron sum-

marizes earlier work on the alumina-water system. According to Rooksby, gibbsite

(AI(OH)3) is the only naturalll' occurring pol1'morph of aluminum hydroxide as, at that

time bayerite (AI(OH)3) and the unusual form nordstrandite (Al(OH)a) were known only as

products of synthesis. A year later, Hathaway and Scblanger (1962) reported nordstrandite

as a mineral from Guam, and Wall and others (1962) from Sarawak (Borneo). And in

1963, Bentor and others described the first natural occurrence of bayerite corroborated by

X-ray difiraction. Regarding conditions of synthesis, Rooksby stated that the slow acidifi-

cation with COz of a sodium-aluminate solution results in precipitation of aluminum

hydroxide with the gibbsite crystal structure iI carried out at temperatures near 100'C and

results in the bayerite clystal structure if the temperatule is nearer room temperature.

Agrng of this bayerite, however, may convert it to gibbsite. He believed that nordstrandite

STRUCTURES OF ALUMINUM HYDROXIDE 45

precipitated from solutions of extremely high pH after long-term aging.In a brief review of the literature, Deer and others (1962) state that gibbsite forms in a

sodium- or potassium-aluminate solution at 30-75"C by autoprecipitation or by introducingpreheated COg into the solution. Bayerite, they stated, forms by more rapid precipitationfrom cool or less alkaline or even acid solutions.

The first synthesis of nordstrandite by Van Nordstrand and others (1956) yieldedmixtures of all three polymorphs in solutions whose pH ranged from 7.5 to g. Subsequentwork by Papde and others (1958) showed that the proportion of nordstrandite increased ifthe pH of the solution was high (pH 13 in one case).

Summing up the results of earlier work, Bye and Robinson (1964) proposed a schemeof aging whereby amorphous aluminum hydroxide gel crystallizes to pseudoboehmite whichin turn converts to bayerite and finally to gibbsite. Bye and Robinson stated that thesechanges are spontaneous at temperatures below 70'C in the presence of the mother liquor,and the rate of each conversion increases with increasing pH and decreases in the presenceof foreign anions. In their own slmthesis ri/ork, Bye and Robinson used aluminum s-butox-ide and various organic liquids as starting ma.terials in order to avoid the obstructive efiectsof foreign ions. They obtained pseudoboehmite precipitates that converted in about 10-20hours to bayerite, but did not detect gibbsite even after aging for 1 year.

Hauschild (1963) managed to produce pure nordstrandite by various techniques, allof which employed aqueous solutions of ethylenediamine for aging the aluminum hydroxidegels. Presumably, the large organic anions have less efiect in hindering crystallization.

Hsu (1964) investigated the complications caused by foreign anions and found that theinhibition is directly proportional to concentration of the anion, and the efficiency is in theincreasing order: chloride(sulfate(phosphate{fluoride. Recent work by Hem (1968) onthe stability of aqueous complexes of aluminum with fluoride, sulfate, and hydroxyl ionsindicates that complexing is partly responsible for the efiects observed by Hsu. Complex-ing of aluminum by anions permanently reduces the amount of aluminum available forprecipitation. In addition, there apparently is a slowing of crystallization due to entrap-ment or proxying of other anions for hydroxyl ions. This effect can be overcome by al-lowing suficient time for crystallization (John D. Hem, oral commun., 1969).

Experiments by Herbillon and Gastuche (1962) using aluminum chloride, nitrate, andsulfate showed that only amorphous gel or pesudoboehmite would form regardless of pH.They ascribed this to ionic impurities which they then removed by dialysis. This per-mitted the relatively rapid precipitation of pure gibbsite at pH 4.6, mixed bayerite andpseudoboehmite at pH 6.5, and mixed gibbsite and bayerite at pH 8.

A study by Barnhisel and Rich (1965) indicates that, at room temperature, gibbsiteforms in acid solutions, nordstrandite forms in slightly acid to neutral solutions, andbayerite forms in alkaline solutions. Hsu (1966) and Hem and Roberson (1967) reportedsimilar results in regard to the acid and alkaline environments favoring precipitation ofgibbsite and bayerite, respectively. Hsu (1966) proposed that the difierence between theformation of the two polymorphs was in the source of the hydroxyl ion. When a stronglyalkaline solution is used, the hydroxyl ions almost immediately couple with the aluminumions to form bayerite. But in acid solutions the hydroxyl ions must come from ttre slowdissociation of water wittr the consequent formation of gibbsite.

Hsu and Bates (1964a) believe that in acid solutions the aluminum ion reacts withhydroxyl ions, as they become available, according to the following steps:

A13+ + (OH)l_ : Al(Ot1;z+

Al(oH)'z+ + (oH)'-: Al(oH)l+

6Ar(oH)l+: Ar6(oH)il.

46 ROBERT SCHOEN AND CHARLES E. ROBERSON

This last ionic species, AI6(OH)IzF, constitutes a hexagonal closed ring. Further polymeriza-

tion can yield higher polymers such as A!n(OH)zz8+ or Alrs(OH)36e+. But, as the positive

charge on the poly'rner increases, further poly'rner'uation slows down because of eiectrical

repulsion. Only when sufficient hydroxyl ions form to discharge the polymers, will they

cluster and form crystalline AI(OH)a.Hem and Roberson (1967) proposed a somewhat difierent mechanism to provide an

explanation of the observed behavior of aluminum hydroxide precipitates at all pH values.

At pH values two or more units below the point of minimum aluminum solubility (about

pH 5.8), the aluminum ion is coordinated by six polar water molecules. The strong positive

charge of the aluminum ion repels the two protons of each water molecule and, as pH in-

creases, eventually drives off a proton forming the monomeric complex ion AI(OH) (OHtb'+.

At about pH 5, this complex ion and hydrated aluminum are in equal abundance with

more of the hydrated aluminum at lower pH values and more of the complex ions at higher

pH values. Two of these complex ions may join by losing two water molecules to form a

dimer:

2A1(OH) (OHt!* : er,1On),(oH,);+ + 2H,O.

Furttrer deprotonation, dehydration, and polymerization of monomers and dimers yields a

ring structure of six octahedrally coordinated aluminum ions with the formula:

A16(OII)rr(OHr)16l'

Coalescence of rings into layers by further growth and stacking of layers results in the

formation of gibbsite (AI(OH)B). Above the point of minimum solubility, precipitation of

bayerite must involve the anion Al(OH;ot-, generally considered to be the dominant alu-

minum species in alklaine solutions. The reaction may be:

Al(oH)i * H'* : Al(oH)3 + H,o.

Reading this reaction to the right, crystalline Al(OH)s (bayerite) precipitates u'ith an ac-

companying increase in pH of the solution.This summary of some of the previous work on the synthesis of aluminum hydroxides

shows the confusion that exists, especially in the earlier work, regarding the environment

in which the various solids precipitate as well as the identity of these solids. As Hsu (1967)

points out, conflict and confusion in the literature on the alumina-water system stems from

the metastability of intermediate products and siow reversibility of reactions. In addition,

we would mention tle problem of incomplete identification of solid products. This paper is

concerned primarily with clarifying the identification of synthetic aluminum hydroxide

precipitates.

Mnrnon or SYNTTTESTs

We prepared most of the synthetic aluminum hydroxides described in this report by

mixing three aqueous solutions. Solution I contained ions of hydrogen, perchlorate, sodium,

and aluminum (4.53X10-4 molar). Solution II contained ions of sodium, perchlorate, and

hydroxyl (6.53X10-s molar). Solution III contained ions of sodium, perchlorate, and alu-

minum (9.06X 10-a molar). The desired amount of solution II (20.0 to 36.5 ml) was added to

25.0 ml of solution I in a polyethylene bottle with vigorous stirring. An amount of solution

III equal in volume to solution II was added at the same time to maintain the concentra-

tion of aluminum at 4.53X10-a molar. The rate of addition was about 1 ml per sec. We

chose the perchlorate anion for this study because of its freedom from a tendency to com-

plex with aluminum. The perchlorate concentration was 0.01 molar in all solutions in order


to maintain a relatively constant low ionic strength (and, therefore, a constant activitycoefficient for reacting ions).

The mixed solutions were stored in a COz-free atmosphere at2soc+2o for periods rang-ing from 16 hours to 10 days. At the conclusion of aging, we measured the pH of the solu-tions with glass and calomel electrodes and then filtered the solutions through plastic mem-brane (Millipore) filters having pore diameters of 0.45 pm (in a few experiments 0.10 pmfilters were used). Immediately after filtration, we affixed the piastic membrane filter to aglass slide with rubber cement and X-rayed it with a Norelco or Picker diffractometer andcopper Ka radiation.

The amount of precipitate often was small (-5 mg) and the X-ray diffractograms were,therefore, incomplete. In order to make positive identification, we prepared larger amountsof solution which yielded as much as 50 mg of precipitate. This was then filtered, dried atroom temperature, scraped off the filter, ground lightly, and X-rayed in a standard alu-minum holder to provide a complete X-ray difiractogram with reduced preferred orienta-tion. Samples for electron microscopy were prepared in a standard manner (Kay, 1965) byallowing a crop of mother liquor, containing some precipitate, to dry on Formvar plasticfilm and using carbon shadowing.

We made all our measurement of X-ray peak intensity on peak height above back-ground rather than on total area. Where the half-widths of X-ray peaks difiered consider-ably, we made no attempt to compare reported intensities quantitatively.

We did not chemically analyze any of the aluminum hydroxide precipitates describedin this report. The matching of X-ray diffraction peak positions and intensities betweenstandards and precipitates was considered adequate for identification. We are aware ofreports proposing that small amounts (usually less than one percent) of sodium and potas-sium found in aluminum hydroxide are necessary to stabilize the lattice (Saalfeld and others,1968). The crystallochemical dificulties of incorporating sodium or potassium ions into thealuminum hydroxide lattice, and the kinetic efiects of other foreign ions, mentioned else-where in this paper, lead us to believe that purity of precipitate is cf no consequence to thisstudy.

Rrsur-rs AND INTERPRETATToN

Table 1 lists the results of the X-ray study of the synthetic aluminumhydroxides. They were prepared under a variety of conditions, and thesample identifier listed in the first column of Table 1 can be used todetermine what these conditions were. The starting pH (related to theamount of hydroxyl-containing solution II added) varied from acid insamples with Arabic numeral 1 to alkaline in samples with Arabic nu-meral 5. The Ietters A, B, and C in the sample identifi.er refer to agingperiods of 16 hours, 2 days, and 10 days, respectively. In order to assessthe reproducibility of the results, we prepared a second group of sampleswith almost identical starting pH values. Roman numeral II identifiesthis second group of samples, and Roman numeral I the first group.

Column 2 of Table 1 lists the ratio of the hydroxyl ions bound toaluminum, to the total aluminum present. This represents an averagefor the number of hydroxyl ions bound to each aluminum ion in thesystem, including both aluminum in solution and in the precipitate. Aratio of 3.00 indicates that just enough hydroxyl ions were added to

48 ROBERT SCHO]iN AND CIIARLES E. ROBERSON

form AI(OH)a with ail the aluminum present. A ratio less than 3.00 indi-cates an acid solution in which some aluminum remains in solution as ahydrated ion or a cationic complex, and a ratio above 3.00 indicates analkaline solution in which some aluminum remains in solution as theanionic complex Al(Ott;nt-. Column three lists the length of aging andcolumn four, the final pH of the solution before filtering. The last sixcolumns of Table 1 list the equivalent interplanar spacing in angstromsand intensity in counts per second of three distinct gro^ups of X-raydiffraction peaks. These groups are 4.7 to 4.9,4.4, and 2.2 A.

Table 1 shows that the number of X-ray peaks and their intensitygenerally increases with increasing final pH of the mother liquor andwith increasing aging time. The 4.8 A peak is developed best, however, inthose intermediate pH samples (3C1) aged the longest (Fig. 1A). Pro-longed aging of- the high pH samples (for example 4AI, 4BI, and 4CT)causes the 4.8 Apeak to disappear (Fig. 18, C, D) as the4.4and2.2 A'peaks increase in intensity. This indicates that the phase that producesthe 4.8 A,peak is different from the phase or phases responsible for the 4.4and 2.2 A peaks, and that these last two peaks probably are caused bythe same phase.

We sought identification of the phases responsible for these threeX-ray peaks from among the three aluminum hydroxide polymorphs:gibbsite, nordstrandite, and bayerite. Figure 1, E and F depict diffracto-grams of pure synthetic gibbsite and bayerite, respectively. Figure 1Gis a diffractogram of natural nordstrandite from Guam (Hathaway andSchlanger, 1965).

Occurrence oI the 2.2 A peak with the 4.4 A peak is characteristic ofbayerite. The reason for the absence of the strong 4.7 1 A bayerite peakand the low intensity oI the 2.2 A peak, as listed in Table 1, will be ex-plained presently. The 4.8 A peak is produced by gibbsite. Verificationof this identification required the study of a larger amount of precipitate(about 50 mg) over an aging period of 76 days. With time, the less intensepeaks of gibbsite appeared and the 4.8 A peak became sharper. Thesample aged for 76 days yielded the complete gibbsite diffractogramshown in Figure 1 E.

Table 1 shows that there is a change in fi.nal pH with aging. This pHdrift is, in general, toward lower values for samples whose initial pH wasbelow the point of minimum solubility (about pH 5.8) and toward highervalues for samples whose initial pH was above that point. Barnhisel andRich (1965, Table 3) observed the same change in pH.

Drift of pH toward Iower values may be due to growth of the gibbsitecrystals during aging. Hem and Roberson (1967 , p. A50) suggested thatthe hydroxyl ions in the gibbsite structure that are on the surface of the

STRUCTURES OF ALU A,IINUM HYDROXIDE

Fro. 1. X-ray difiractograms of tlpical synthetic products developed during short-termaging in tle alumina-water system (A, B, C, D) and Al(OH)s standards (E, F, G).

49

lBcq

iq

R,ffiiII

I

ElII

FiI

iclrlu 112 ?

eq rees '[

2


Terr.r 1. X-ne.v DrrlnectoN Rrsutts ox SvNrrrnnc AI(OH)3

Sampleldentifier

1 A IB IC I

Age

16 hrs2 d,ays

10 days

l'inalpH

4 . 7 34 . 6 64 . 5 6

Spacing and Intensity

d G ) I d ( b r o . j 0 ,

1 A I IB I IC I I

2 . 3 22 . 3 22 . 3 2

16 hrs2 days

10 da1'g

4 . 9 44 . 8 34 . 7 0

4 .70 20 4 .37 24

2 I t 7B IC I

2 . 7 92 . 7 92 7 9

16 hrs2 days

10 days

5 . 2 85 . 1 74 . 8 3 4.90 28 4 .40 24

2 A I IB I IC I I

2 . 7 92 . 7 92 . 7 9

16 hrs2 d,ays

10 days

< a A

5 . 1 05 . 0 2 4.82 20

3 A IB IC I

2 . 9 02 . 9 02 . 9 0

16 hrs2 days

10 days

6 . 1 55 .936. 95 4.88 130 4 .37 60

3 . A ' I IB I IC I I

2.902 .902 .90

16 hrs2 days

10 days

5 .8sq 6 0

6.45 4 . 8 7 3 6 4 . N 2 0

3C II-1 2 . 9 0 10 days 6 .+5 4.90 76 4 .42 52

4 A IB IC I

3 . 1 53 . 1 5J . I J

16 hrs2 days

10 days

8 8 79 . 0 18 .95

4.77 28 4 .40 80 2 .23 244 . 7 7 1 4 4 . 3 6 1 8 8 2 . 2 2 3 6

4.M 160 2 .2+ 28

4 A I IB I IC I I

3 .003 .003.00

16 hrs2 days

10 days

6 . 3 76.498.48 4.42 r+0 2.23 M

5 A IB IC I

3 . 5 2? ( ,

3 . 5 2

16 hrs2 days

10 days

9 . 4 89 . 6 99. 54

4.7s 20 4 37 I20 2.23 364.77 16 4.37 24O+ 2.23 48

+.37 280+ 2.22 48

5 A I IB I IC I I

3 . 5 2s . 5 23 . 5 2

16 hrs2 days

10 days

9 . 2 39 . 4 79.64

4 .37 48 2 .23 164.40 r52 2.23 324.37 232+ 2.22 40

STRUCTURES OF ALUMINUM HVDROXIDE 51

crystal may have an alkaline effect on the pH of the solution. Becausesurface area per unit volume decreases during crystal growth, hydroxylions in the crystals should exert a decreasing effect on the pH of thesolution as the crystals grow larger. Other Iikely explanations for theobserved downward pH drift are the deprotonation reaction duringpolymer growth as well as the dissociation of water proposed by Hsu(1e66).

The preceding hypotheses, however, do not explain the upward pHdrift o{ solutions containing bayerite. Hydroxyl ions on the surface ofbayerite crystals should have little efiect on the pH because the crystalsgrow rapidly in the alkaline environment and quickly achieve large size.X-ray diffraction and electron microscopy indicate that bayerite crys-tallites attain sizes in excess of 1,000 A within 24 hours. Instead, theobserved upward drift in pH probably is due to the reaction:

Al(oH)3 + Al(oH)l- : 2A1(OH)3 + (OH)'-

in which aluminate ions are added to bayerite crystals as slow growthcontinues during aging.

All samples produced some very broad X-ray peaks. There was atendency for a broad peak at about 9 A to migrate toward 7 A with in-creasing crystallinity of the sample. This broad maximum may indicatethe incipient formation of pseudoboehmite, a monohydrate of aluminum.Another broad peak near 4 A tended to migrate toward 3 A in morecrystalline samples, and may also be a pseudoboehmite peak.

Filtration of sample 3CII through a 0.45 trrm membrane produced anopalescent filtrate so we refiltered the filtrate through a 0.10 pm plasticmembrane filter. Listed in Table 1 under identifier 3CII-1 are the resultsof the X-ray difiraction study of this 0.45 pm-0.10 pm size material.Comparison with the data l isted for sample 3CII indicates thatmuchcrystalline material was less than 0.45 pm and therefore was missed inthe study of other samples listed in Table 1. Subsequent studies using0.10 pm plastic membrane filters, however, did not result in any changesin the identification of phases.

Figure 2 depicts the results of one of the later studies using 0.10 pmplastic membrane fi.lters. Samples prepared at (OH)s to Al ratios oI 2.7,2.8, and 2.9, and aged for periods of 7 days, 7 days, and 8 days, respec-tively, gave final pH values of the mother l iquors ol 4.74,4.80, and 4.86,respectively. Figure 2, A and B show no crystalline material by X-raydiffraction. Figure 2C, ol material prepared with an (OU)u to Al ratioof 2.9, exhibits a broad peak at 4.8 A identified as gibbsite by otherlong-term aging experiments. X-ray oI the 2.9 ratio sample t hour afterfiltration (Figure 2D) shows a large increase in intensity of the broad


A

IIIII

2 -7 f resh I

D e g r e e s 2 e ( C u K c )

FIc. 2. X-ray diffractograms showing the development of synthetic gibbsite during aging

and drying as a function of tle ratio (OH)B/AI.


gibbsite peak. X-ray ol the 2.7 ,2.8, and 2.9 ratio samples after 2 days,2 days, and 1 day, respectively, of drying on the plastic membrane filtersresulted in Figure 2, E, F, and G respectively. These diffractogramsindicate that gibbsite formed in all solutions though drying was neces-sary to make the gibbsite in the 2.7 and 2.8 ratio samples detectable byX-ray diffraction. Apparently, many low pH solutions prepared earliercontained gibbsite that was too fine to be trapped on 0.45 prm filters.This may account for the erratic distribution of gibbsite peaks reportedin the upper half of Table 1.

The diffractograms presented in Figure 2 reveal several other thingsabout the gibbsite formed. Diffraction from the (002) or basal planes ofthe gibbsite crystals causes the single X-ray peak. The great breadth ofthis peak indicates that the crystallites are relatively small in the basal(crystallographic c-axis) direction. Calculation of this dimension usingthe Scherrer formula (Klug and Alexander, 1954, p. 512) results in 45 Afor samples 2.7 and, 2.8, and 80 A for sample 2.9. Measurements ofcrystallite shadows on electron micrographs corroborate these dimensions(Ross Smith, oral commun., 1968). But at least one dimension of thesecrystallites must exceed 1,000 A because they were trapped on 0.1 pmfilters. Figure 3 shows a hypothetical gibbsite crystal possessing thesedimensions and the hexagonal form seen in electronmicrographs such asFigure 4. Apparently, the growth habit of gibbsite favors the formationof thin platelets, whose upper and lower surfaces are parallel to thebasal planes responsible for the observed X-ray peak. According to VanNordstrand and others (1956, p. 7t2), ". . . gibbsite usually grows asplatelets parallel to the layers of its layer lattice." (basal planes).

Ftc.3. Sketch of hypothetical gibbsite crystal formed during synthesis.

54 ROBERT SCHOEN AND CHARLES E, ROBERSON

Frc. 4. Electronmicrograph of synthetic gibbsite crystals.

Figure 2, C and D show that drying of the precipitate produced morecrystallites capable of diffracting X-rays. But drying did not appreciablyincrease the size of these crystallites parallel to the c-axis as shown by thelack of change in peak breadth. Perhaps evaporation of interstitialmoisture precipitated more gibbsite crystallites or caused the plateletsto grow only in the o and b crystallographic directions. Removal ofmoisture films from the crystallites would also decrease absorption ofX-ray energy, thus tending to increase the X-ray signal received fromthe crystalline material.

In general, aging of gibbsite or bayerite in the mother liquor results inan increase in the size of the crystallites (Figure 2 versus Figure 1E). Inaddition, a complete development of all crystal planes by aging resultsin all permissable X-ray diffractions being observed as in Figure 1, E andF. The only example of crystallites becoming smaller during aging is thegradual disappearance of gibbsite formed in mixtures with bayerite atintermediate pH values (Table 1). This implies that gibbsite is unstablewith respect to bayerite in these solutions. The rise in pH of the solution,caused by precipitation and growth of large bayerite crystallites, maybe responsible for making the smaller gibbsite crystallites unstable.

The difiractogram of well-crystallized bayerite (Figure 1 F) showsthree peaks of extreme intensity at 4.7 I A, +.S5 A, and 2.22 A. Thenewly-formed bayerite samples of Table 1, however, show only a strong4.4 A peak and a weak 2.2 Lpeak;the4.7l A peak is missing. Part ofthe reason for these anomalous intensities is the rate of development of

STRUCTURES OF ALUMINUM IIYDROXIDE 55

the crystal planes responsible for these three difiractions. The 4.35 Apeak originates from (100) prism planes in the crystal and the 2.22 hpeak orginates from (111) pyramid planes. The 4.7 1A peak originatesfrom (001) basal planes comparable to those responsible for the strong4.85 A peak in gibbsite. Absence of rhe 4.71 A (001) peak in newly-formed bayerite implies a poor development of basal planes in the crys-tallites. Crystallites with well-developed (100) prism faces and less well-developed (111) pyramid faces might look like Figure 5. This agreeswith the reported habit of bayerite as .'. . . rods or tapered rods the longdirection of which appears roughly perpendicular to the layers of thelayer lattice." (basal planes) (Van Nordstrand and others, 1956, p. 714).Results of electron diffraction by Lippens (1961, p. 69-70) support thisperpendicularity between the basal layers and the longitudinal directionof the bayerite particles. An electronmicrograph of bayerite aged for 10days (Fig. 6) proves the general correctness of this concept of the shapeof newly-formed bayerite crystallites.

Frc. 5. Sketch of hypothetical bayerite crystal formed during synthesis.

JO ROBERT SCHOEN AND CEARLES D. ROBERSON

Frc. 6. Electronmicrograph of synthetic bayerite crystals,probably including a few nordstrandite crystals.

There are at least three reasons why the intensities of X-ray peaks

from newly-formed bayerite should not match those of aged bayerite.

First, the particle shape, which is elongated parallel to the c-axis, tends

to produce only (hk0) diffractions because the basal planes at right an-

gles to the c-axis are small initially. But after ten days aging (Fig. 6), the

basal planes are large enough to produce sharp X-ray diffractions. Ab-

sence of these peaks in our 10-day old samples must be a consequence of

the elongated particle shape and the resultant preferred orientation oc-

curring during filtration. Basal planes of bayerite crystallites on the plas-

tic membrane fiIters, though large enough' are not in a favorable position

to diffract X-rays. Figure 1F, the difiractogram of bayerite scraped from

the filter and mounted to minimize preferred orientation, exhibits all

peaks with proper intensities. Finally, the small amount of aluminum

hydroxide precipitate deposited on the plastic membrane f,lter is not

sufficient to satisfy the X-ray requirements for a sample of infinite thick-

ness (Alexander and Klug, 1948). X-raypeaks at high two-theta angles,

therefore, are especially subject to diminution of intensity with respect

to lower angle peaks. This explains the relative weakness of the 2.2 A

peak.The growth of gibbsite in acid solution and bayerite in alkaline solution

STRACTURES OF ALUMINUM HYDROXIDE 57

occurs by the addition of units of aluminum hydroxide. These unitsattach themselves at a greater rate to the edges of gibbsite crystallitesthan to the top and bottom surfaces. The result is that the slowest grow-ing faces (that is, the top and bottom surfaces) eventually form theprominent crystal faces of gibbsite. Conversely, the aluminum hydroxideunits attach themselves preferentially to the basal planes of bayeritecrystallites with the result that the slower growing side faces (prism andpyramid) eventually form the dominant crystal faces.

To complete our understanding of the crystal form of the three alu-minum hydroxide polymorphs, we must examine the electron microscopeand electron diffraction data of Lippens (1961, p. 73-74) on nordstrand-ite. He, as well as Hauschild (1963, Fig. 2) found that nordstranditecrystallites take the form of long rectangles or more rarely long parallelo-grams. The few small rectangles and parallelograms in Figure 6 corre-spond closely to those figured by Lippens and Hauschild and may benordstrandite.

The cross section of nordstrandite crystallites is hexagonal, and, as inbayerite, the c-axis parallels the long direction of the crystallites. Butthe D-axis always appears to coincide with the width of the crystallitesindicating a relatively poorer development of the o-axis direction. Lip-pens found the o-axis direction to range between one-half and one-sixththe length of the D-axis direction. Figure 7 is a sketch of a hypotheticalnordstrandite crystallite embodying these characteristics.

b - a x i s

Frc. 7. Sketch of hlpothetical nord-strandite crystal formed during s1'nthesis.

c - a x t s


None of the samples described thus far showed X-ray evidence of con-taining nordstrandite, the third polymorph of aluminum hydroxide. Asalready described, much confusion exists in the literature regardingconditions of synthesis of nordstrandite. Ideal pH conditions seem torange from "slightly acid to neutral" (Barnhisel and Rich, 1965) to "pH13" (Pap6e and others, 1958).

To help resolve these inconsistencies, or at least to provide additionaldata, we prepared a series of four solutions for long-term aging whoseinitial pH ranged from 9.46 to 12.0. Concentrations of ions were greaterin these solutions than in those previously described, and in one solutionpotassium and sulfate were present for use in another experiment.

We had hoped to fi,nd the aluminum hydroxide polymorphs bayeriteand nordstrandite restricted to well-defined pH regions. Or, failing this,we hoped to discern a distinct increase in the amount of one phase anddecrease in the amount of the other during aging. The X-ray results, todate, are equivocal. Pure bayerite without admixed nordstrandite formedinitially in the pH 12.0 solution. But within 90 days aging, X-ray detect-able nordstrandite formed as the pH increased to 12.82. The other threeless alkaline solutions produced bayerite admixed with nordstranditealmost from the start. Nordstrandite X-ray peaks increased in intensityduring two and one-half years of aging, but so did the bayerite peaks.To date, peak ratios show no distinct trend.

In contrast to these results, an electronmicrograph (Fig. 8) of materialsimilar to that shown in Figure 6, but aged for two years, shows a definiteincrease in the size of nordstrandite crystallites relative to bayerite.Figure 8 also shows rounding of the edges of bayerite crystallites andthe development of etch pits parallel to the basal cleavage. These featuresindicate resorption of the bayerite into the mother liquor.

These results imply that slightly alkaline environments promote theprecipitation of nordstrandite and strongly alkaline environments theprecipitation of bayerite. But bayerite is unstable with respect tonordstrandite in alkaline solutions.

Srnucrune on Ar-uurxuu Hvonoxrpe Por-vuonpns

The preceding experimental results corroborate the observations ofBarnhisel and Rich (1965) that acid environments favor the crystalliza-tion of gibbsite, neutral environments nordstrandite, and alkaline en-vironments bayerite. In order to understand why the pH of the aqueousenvironment controls the structure of the aluminum hydroxide poly-morph, we must review what is known about the detailed structures ofthe three phases.

STRUCTURES OF ALUMINUM HYDROXIDE

Frc. 8. Electronmicrograph of synthetic bayerite (barrel-shaped) and nordstrandite (long

sharp-cornered rectangles) aged at about pH 9 for two years.

Gibbsite. According to Megaw (1934), gibbsite is monoclinic (pseudo-hexa-gonal) wi th o:8.624 L, D:5.060 L, c :9.700 A, and 0:94o341. Thesedimensions have remained essentially unchanged by more recent investi-gators, although Saalfeld (1960) described a natural triclinic variety.

The simplest structural unit of gibbsite, a layer, consists of two sheetsof close-packed negatively charged hydroxyl ions bound together bypositively charged aluminum ions that occupy two-thirds of the inter-stices between the hydroxyl sheets (Fig.9). Growth of this layer ofgibbsite occurs by superposition of other layers and by lateral extensionof layers. During the Iateral extension of gibbsite layers, the constraintson ionic bonding are that each aluminum ion be bound to six hydroxylions by bonds of one-half strength, and that each hydroxyl ion be coordi-nated to two aluminum ions by bonds of one-half strength and also coordi-nated to one "hole" or interstice where there is no cation. This satisfi.esthe charges on the aluminum and hydroxyl ions and no ionic forcesremain to attach superposed layers.

59

l'a,l:::::


S et at 1t3;2/3@ at at 1/6;5/6

Fre. 9. Diagrammatic representation of

idealized gibbsite layer structure pro-jected on a plane perpendicular to the o-

axis (from Megaw, 1934).

Incomplete filling of the interstices of the gibbsite layer results in the

"holes" being less electrostatically positive than surrounding portionsof the cation sheet. This asymmetry of charge distribution has an affect

on the orientation of the oxygen-proton bond within each hydroxyl ion.The protons, instead of being oriented straight up away from the cationsheet, tend to be displaced toward the neutral holes.

In addition, the very high polarizing power of the aluminum ion pro-

duces a more signifi.cant distortion of the hydroxyl ions. Bernal andMegaw (1935) proposed that this polarization produces a hydroxyl ionwith tetrahedral symmetry analogous to the water molecule. Two alumi-num ions bind two of the negatively charged corners of the four-corneredion. The remaining negatively charged corner and the positively chargedcorner containing the proton are available for the formation of ionicbonds with other hydroxyl ions, both within the gibbsite layer and be-tween layers. A distribution of both positively and negatively chargedcorners of tetrahedral hydroxyl ions on the surfaces of the gibbsite layersfavors the buildup of superposed gibbsite layers. A bond forms betweenthe negatively charged corner of one hydroxyl ion and the positivelycharged corner of a hydroxyl ion in the superposed layer. There is oneso-called hydroxyl bond for each hydroxyl ion.

TIIII

o\o)l l

G

oI

III

-iL-

STRUCTURES OF ALUMINAM HYDROXIDE 61

The constraints imposed by this type of bonding result in the hy-

droxyl ions of each superposed layer of gibbsite lying directly on top of

the hydroxyl ions of the Iayer below in a configuration resembling open-

packing (Fig.9). The effective radius of hydroxyl ions in one layer to

hydroxyl ions in a superposed layer (that is, B-B or A-A in Fig. 9) is

only 1.39 A, close to the 1.34 A effective radius between hydroxyl ions

and the cations (Table 2). This indicates that distortion of the hydroxyl

Tastn 2. Ellrcrrvn INrsnroNrc DlsraNcns tN Grnrsnr lrro Bnucrrr(from Bernal and Megaw, 1935)

Efiective hydroxylradius toward hydroxyl

in the next layer

Effective hydroxyl

radius toward cation

1 .34 A1 .39 A

1 . 3 e A1 . 6 1 A

ion by polarization produces an "open-packing" in which the centers of

ions are almost as close as in true close-packing. This stacking arrange-

ment of double-sheet gibbsite layers can be described by the symbol:

AB I ee I nn and so forthl l

where A and B denote hydroxyl sheets in different positions as shown in

Figure 8, and the vertical dashes represent the boundaries between

double-sheet layers.Bayeri.te. The absence of large, single crystals of bayerite with which to

make structural determinations, renders the detailed structure of this

polymorph uncertain. Montoro (1942) concluded from X-ray powder

analysis-that bayerite possessed a hexagonal lattice with a:5'01 A,

c:4.76 A, containing 2 molecules of AI(oH)3. This is a layer structure

similar to gibbsite but without the distortion of the sheet of anions

around each vacant cation site that makes gibbsite pseudo-hexagonal.

Instead of the hydroxyl ions in adjacent layers being in a state of open-

packing as in gibbsite (Fig. 9), the hydroxyl ions interdigitate in a state

of close-packing (Fig. 10). This proposed structure for bayerite is identi-

cal with the structure of brucite (Mg(OH)), except that in brucite every

cation position is filled and in bayerite one-third of the cation positions

are vacant.Milligan (1951) and Kroon and Stolpe (1959) took exception to the

iseI


brucite-like structure proposed by Montoro. Other workers (Yamaguchiand Sakamoto, 1958; Lippens, 1961) using material of greater puritythan Montoro's, confirm the general correctness of the brucite model forbayerite. Lippens showed that bayerite is slightly distorted from hexag-onal symmetry and is more l ikely orthorhombic with o: 8.67 A, D:5.06A, and c:4.71A. For comparison, the hexagonal brucite model of Yama-guchi and Sakamoto yields an orthohexagonal o-axis of 8.74 A.

The brucite structure consists of two sheets of close-packed, negativelycharged hydroxyl ions held together by positively charged magnesiumions occupying all the interstices between the hydroxyl sheets. These twosheets of hydroxyl ions, together with the intervening sheet of magnesiumions, constitute a layer of the mineral brucite, and this layer is the smallestunit that can exhibit the propdrties of brucite. Larger brucite crystalsform by the superposition of layers on top of one another, as well as bythe lateral growth of all the sheets. Lateral growth occurs simply by theaddition of magnesium and hydroxyl ions in a configuration where everymagnesium ion binds six hydroxyl ions with bonds of one-third strengthand every hydroxyl ion binds three magnesium ions. These bonds areionic and three, one-third strength magnesium-hydroxyl bonds com-pletely satisfy the charge of a hydroxyl ion. Therefore, there are no ionicforces available to bond together the hydroxyl ions of superposed brucitelayers. In addition, the proton-oxygen bond within each hydroxyl ionpoints straight up at right angles to the sheets of ions and away from thesheet of magnesium cations (Elleman and Williams, 1956). This resultsin a sheet of protons on the upper and lower surfaces of each layer ofbrucite, and it produces a tendency toward cylindrical symmetry in thehydroxyl ions.

Growth of brucite by the superposition of layers, therefore, must takeplace without the help of any ionic bonding and must overcome theelectrostatic repulsion of sheets of adjacent protons. The bonding be-tween brucite layers is due to van der Waals forces supplemented byforces related to the interaction of the rotating hydroxyl dipole ions. Asa result of all these constraints, the hydroxyl ions in successive brucitelayers interdigitate in such a way as to keep adjacent protons as farapart as possible. This interdigitation, similar to that found within thebrucite layer, has been called "close-packing" but it is much more openthan that. Table 2 shows that the effective hydroxyl ion radius towardhydroxyl ions in the next layer in brucite is 1.61 A (t.SZ A in bayeriteaccording to Lippens, 1961), whereas the effective hydroxyl ion radiustoward a cation is 1.39 A. tfrir indicates how much more open this so-called "close-packing" between layers of brucite really is.

This discussion of the structure of brucite (Mg(OH)z) applies directly


to the structure of bayerite (AI(OH)a) if we keep in mind the substitution

of two aluminum ions and one vacancy for every three magnesium ions.

The presence of holes in the cation sheet introduces some distortion into

the bayerite structure compared with the perfect hexagonal symmetry

of brucite. Bayerite, however, does not suffer the additional distortion

caused by hydroxyl bonding within the layers that affects gibbsite'

Comparison of Figures 9 and 10 shows that the simplest unit of struc-

ture of gibbsite and bayerite is identical (a double sheet of hydroxyl

ions), and the difference between the two polymorphs is solely in the

manner of stacking superposed layers. The symbol AB IAB IAB andso forth, describes the stacking scheme for bayerite as depicted inFigure 10.

$""minumFrc. 10. Diagrammatic representation of brucite-bayerite layer structure.

Nordstrandite. Because of the lack of good single-crystals and the relative

recency of the discovery of nordstrandite, both as a synthetic product(Van Nordstrand and others, 1956) and as a naturally occurring mineral(Hathaway and Schlanger, t962; Wall and others, 1962), the detailed

B

AT-

I

+I IO

uB

A

- . / \

ROBERT SCHOEN AND CHARLES E. ROBERSON

structure of nordstrandite remains obscure. By analogy with the X-raypatterns of gibbsite and bayerite, Van Nordstrand and others (1956)proposed that nordstrandite (called "bayerite II" by them) possessed alayered structure with a spacing between layers equal to the average ofthe gibbsite and bayerite spacings. They proposed that the new phasemight be a screw dislocation polymorph in which the gibbsite packing andbayerite packing alternated and in which the Burgers vector was twolayers high.

Lippens (1961) examined the structure of nordstrandite as well asgibbsite and bayerite. His powder samples for X-ray study containedgibbsite, bayerite, pseudoboehmite, and nordstrandite in varying pro-portions, requiring him to ignore X-ray lines attributable to the otherthree phases. For this reason, as well as the lack of single-crystal data,Lippens' results are only a reasonable approximation to the structure ofnordstrandite.

Lippens (1961) reported that his synthetic nordstrandite was mono-clinic with a:8.634, D: S.Ot A, c : 19.12 h, and B : 92.000. The c-dimen-sion of I9.l2L indicates that the unit cell of nordstrandite contains twolayers of gibbsite (4.S5 A each) and two layers of bayerite (4.71 A each).These may be stacked in either of two ways: I-bayerite, gibbsite, gibbsite,bayerite; or ll-bayerite, gibbsite, bayerite, gibbsite. By an analysis ofthe 00/-diffractions, Lippens concluded that arrangement II was mostprobable. This coincides with the arrangement proposed by Van Nord-strand and others (1956), and is depicted in Figure 11. The symbol forthis stacking sequence is:

ABluelt l

An attempt at single-crystal studies on the natural material fromGuam by Hathaway and Schlanger (1965), met with failure when allnordstrandite crystals selected proved to be multicrystalline or dis-torted. A similar attempt on the natural material from Sarawak ap-parently was successful (Saalfeld and Mehrotra, 1966).

As recently reported by Saalfeld and Jarchow (1968), nordstranditeis t r ic l in ic wi th o:8.752 A, b:5.069 L, c : t0.244 A; o:109.33o, B:97 .66", ?:88.340. These authors report that "From the structural pointof view nordstrandite can be regarded as intermediate between gibbsiteand bayerite." But unlike the stacking sequence of nordstrandite pro-posed by Lippens (1961), Saalfeld and Jarchow (1968), in their Figure 3,depict the unusual stacking sequence:

B A I A CI

BAiI

ABiI

and so forth.

CB and so forth

STRUCTT]RES OF ALUMINUM HVDROXIDE

N':o)

il

ccL

' ,lt

I

65

N) a l u m i n u mw

Frc. 11. Diagramatic representation of nordstrandite layer structure showing alternatiorr

of gibbsite and bayerite modes of stacking.


The validity of this model requires further elaboration and confirmation.Pending a complete and precise description of the nordstrandite

structure, the regular interlayer model (Fig. 11) has much in its favor.By analogy with another large group of layer structures, the clay min-erals, interlayering is a common phenomenon in response to environ-mental changes. And, as in the clay mineral group, there seems to be adefinite tendency for a 1i1 regular interlayered structure to be morestable than either irregular interlayering or other proportions of thecomponent layers. In addition, a single layer of aluminum hydroxidepossesses a marked proclivity to interlayer, both in nature (the chloritemineral group) and with other clay minerals in the laboratory (Hsu andBa tes ,1964b ) .

Sranrr,rrv or AlutrrNulr HynnoxrnB Porvuonpgs

Several suggestions have been offered to explain why gibbsite or bayer-ite may be favored during precipitation of aluminum hydroxide. VanNordstrand and others (1956) suggested that if nordstrandite is a screwdislocation polymorph with a Burgers vector of two layers, bayerite canbe considered to be a screw dislocation ploymorph in which the Burgersvector is one layer. They suggested that unspecified constraints imposedby this type of growth might prevent the gibbsite-type open-packingfrom developing. Barnhisel and Rich (1965) believed that some of theirdata showed a "seeding" effect by foreign mineral surfaces that mightcontrol the structure of the aluminum hydroxide polymorph. Hsu (1966),realizing the great difference in the rate of precipitation of gibbsite andbayerite, suggested that rapid precipitation may favor the bayeritestructure, and slow precipitation the gibbsite structure. None of thesesuggestions, however, specified the mechanism of stability control.

Our study leads us to conclude that a difierent control from thosealready proposed is responsible for the.structures of the aluminum hy-droxide polymorphs. Comparison of Figures 9 and 10 shows that the onlydifference between gibbsite and bayerite is in the mode of stacking thealuminum hydroxide layers. And the two different modes of layer stack-ing reflect the shapes of the hydroxyl ions on the opposing sheet surfaces.If mild polarizing forces affect the hydroxyl ions, they assume a cylindri-cal shape, and the proton-containing ends of the ions move as far aspossible away from the aluminum ions between the sheets of hydroxylions. When layers of aluminum hydroxide with this hydroxyl ion con-figuration superpose, the repulsive forces of the opposing protons forcethe layers to interdigitate. The brucite-bayerite type structure results.When stronger polarizing forces affect the hydroxyl ions, the cylindricalsymmetry breaks down and tetrahedral ions form. The distribution of


electrostatic charges on these tetrahedral ions permits an ionic bond toform with tetrahedral hydroxyl ions in the superposed aluminum hy-droxide layer. The strong attraction of these ionic hydroxyl bonds re-quires that the hydroxyl ions arrange themselves vertically above oneanother perpendicular to the layers. The gibbsite structure is the result.It is probable then that formation of gibbsite or bayerite depends uponthe distortion, and consequently the degree of polarization, of the hy-droxyl ions.

But how does the pH of the mother liquor cause such a difference inpolarization of the hydroxyl ions in the precipitate? And what can ac-count for the maintainence of a large disparity in degree of hydroxyl ionpolarization in two solid phases of the same composition? The answersto these questions constitute what is, at present, only a reasonable guess.Additional specific investigations must be undertaken to evaluate thefollowing hypothesis.

First, we must remember that aluminum is amphoteric and can existin solution as a complex cation or anion depending on the pH. Proofof this statement is afforded by the observation that electrolysis of asolution of potassium aluminate causes deposition of aluminum hy-droxide on the anode, whereas electrolysis of a solution of aluminumnitrate deposits aluminum hydroxide on the cathode (Berges, 1947). Asdescribed previously, Hem and Roberson (1967) showed that the pre-dominant cationic aluminum complex between pH 4 and pH 5.8 isAl(OH)(OHr)r'+. In this complex ion, the lull polarizing power of thealuminum ion affects a single hydroxyl ion. Polymerizatiort of the mono-meric AI(OH)(OH)rrt+ brings the polarizing power of additional alumi-num ions to bear, and deprotonation of the water molecules continueswith the formation of strongly polarized hydroxyl ions. The precedingchain of events would be expected in acidic solutions. Above pH 5.8,however, the dominant aqueous aluminum species is the complex alumi-nate anion AI(OH)41-. We would expect the excess negative charge tonullify the polarizing power of the aluminum ion and form cylindricalhydroxyl ions. Therefore, the phase that precipitates from acid solutionsshould contain highly polarized hydroxyl ions because of their abundancein the acidic mother liquor, and the phase that precipitates from alkalinesolutions should contain weakly polarized hydroxyl ions because of theirabundance in the alkaline mother liquor. And this is exactly what weobserve.

Judging from the preponderance of gibbsite in nature, we can inferthat strong polarization of the hydroxyl ion in aluminum'hydroxide isthe most stable state under conditions prevailing at the earth's surface.It terms of crystal energy (Pauling, 1960, p. 509), the longer hydroxyl-

68 ROBBRT SCHOEN AND CHARLES E. ROBERSON

hydroxyl bond in bayerite, caused by the weak polarization of the hy-droxyl ions, creates a smaller crystal energy than strongly polarizedhydroxyl ions do and, therefore, a less stable structure. Bayerite shouldtend to undergo a slow increase in the polarization of its hydroxyl ionsand recrystalize to gibbsite.

Our experimental results, however, do not support this. What formsduring long-term aging of bayerite is not gibbsite but nordstrandite, amineral of intermediate structure containing both strongly and weaklypolarized hydroxyl ions. Nordstrandite, which is probably a regular inter-layering of the gibbsite and bayerite modes of stacking, may be the stablepolymorph of aluminum hydroxide in an alkaline environment.

A probable qualitative explanation for this is that even though thecrystal energy associated with weakly polarized hydroxyl ions is lessthan that associated with strongly polarized hydroxyl ions, the relativeactivity of strongly polarized versus weakly polarized hydroxyl ions inthe solution also afiects the stability of the solid phase. If the solutionis alkaline, the activity of weakly polarized hydroxyl ions (AI(OH)a1-)is much greater than that of strongly polarized hydroxyl ions (AI(OH)(OH)r)6'+)(Hem and Roberson, 1967). This ratio of ion activities maymake a structure containing both strongly and weakly polarized hy-droxyl ions (nordstrandite) more stable than one containing only stronglypolarized hydroxyl ions (gibbsite).

In other words, strongly polarized hydroxyl ions make a crystal cagethat is more efficient at containing the internal energy of the enclosedatoms. But this crystal cage must be in a state of dynamic equilibriumwith the solution surrounding it. If the surrounding solution contains anabundance of weakly polarized hydroxyl ions, the crystal cage mustcompensate for this. It does this by assuming a structure made up ofstrongly polarized hydroxyl ions for high efficiency in containing theinternal energy of its atoms, and embodying weakly polarized hydroxylions in response to the high relative activity of these ions in the surround-ing environment.

GBocsnursrnv or AluurNUM HyDRoxrDE MTNERALS

We now can compare the laboratory observations and hypothesis forpolymorph stability with the natural occurrence of aluminum hydroxideminerals to see if we can obtain additional eeochemical insieht.

Gibbsite in nature. The most abundant aluminum hydroxide mineral isgibbsite. It is the principal mineral in most bauxite ore deposits, theprimary source of aluminum metal. Gibbsite is also abundant in laterite,the ferruginous analog of bauxite. Harder (1952) described the geology

STRUCTURES OF ALAMINUM HYDROXIDE 69

of many bauxite deposits and ascribed their genesis to the interplay ofmany factors.

Bauxite (and its dominant mineral gibbsite) is principally a residualweathering product of any kind of aluminum-containing rock. There is acomplete spectrum of parent rock types from alkali-rich, aluminousigneous rocks through clays, basalts, and metamorphic rocks to alumi-num-poor limestones. In addition to an easily soluble aluminous rock,warm climate, high rainfall, moderate topographic relief, and freedomfrom erosion are all necessary for the optimal development of bauxitedeposits. The advent of extensive erosion will, of course, remove bauxitefaster than it f orms. And a change in the delicate balance of other factorswill alter the geochemical conditions and end the formation of bauxite.

Some believe that bauxite deposits form by solution, transport, andprecipitation of aluminum. Theobold and others (1963) summarized thebackground of this hypothesis and presented a small-scale natural ex-ample. They postulated the transport of aluminum in acid solution and(or) as a sulfate complex, derived from the oxidation of disseminatedpyrite in the parent rock.

Many bauxite deposits of unquestioned residual origin show evidenceof local solution and precipitation. It is still uncertain, however, whetherthis mechanism can produce large bauxite deposits some distance awayfrom the source of aluminum.

One type of bauxite deposit, whose genesis still is an enigma, developson aluminum-poor limestones. It seems quantitatively impossible toconcentrate the one tenth of one percent or less of Al2O3 in the insoluble

residue of these limestones to form thick bauxite deposits. Perhaps theseunusual bauxites represent accumulations of aluminous material derivedfrom a wide area and washed into sinkholes and other Iocal basins. Dur-ing filling of a sinkhole, large amounts of water percolate through thedeposited sediments. And, after filling, peat bogs frequently develop atthe surface resulting in continued percolation through the sediments ofwaters rich in humic acid (Clark, 1966; Bushinsky, 1964).

In addition to gibbsite, boemite (AIO(OH), one of the monohydrateminerals of aluminum, frequently occurs in tropical bauxite deposits nearthe surface where the heat of the sun may cause dehydration of gibbsite

(Harder, 1952, p. 56). Although the temperature for this transitionranges from 1300C to 200oC in the laboratory, the presence of salts in thesoil profile may reduce the activity of water suficiently to make this

dehydration possible at lower temperatures. Hsu (1967) showed thathigh salt concentrations would delay almost indefinitely the conversion

of pseudoboehmite to bayerite. Boehmite also is the principal aluminousmineral in bauxite deposits around the Mediterranean Sea where regional

70 ROBERT SCHOEN AND CHARLES E, ROBERSON

metamorphism of the rocks may have caused dehydration. Similarly,Hill (1955) noted abundant boehmite associated with faults in Jamaicanbauxite and concluded that the dehydration of gibbsite was in responseto the mechanical breaking of the rocks.

Bauxite frequently contains kaolinite which reduces the grade of thealuminum ore. Kaolinite generally forms by the leaching of other alumi-nosilicate minerals and is a precursor in the formation of gibbsite. Italso may be a product of the resilication of gibbsite. Iron oxides such ashematite and goethite vary in amount from a few percent in bauxites upto several tens of percent in laterites. Resistant minerals from the parentrock such as quartz and anatase often become concentrated in the resid-ual bauxite.

All the geologic and mineralogic evidence presented points to theformation of gibbsite in an environment with a rapid through-flow ofmeteoric water, where kaolinite forms from primary aluminosilicates, andwhere quartz is preserved. An environment of moderate acidity satisfiesthese several requirements. And this is in complete accord with ourlaboratory results where the synthesis of gibbsite was restricted to acidsolutions.

Nord.strandite in Nature. Firstrecognized as a synthetic product byVanNordstrand and others, (1956), Hathaway and Schlanger (1962; 1965)discovered natural nordstrandite on the island of Guam, and Wall andothers (1962) on the island of Borneo. Careful reexamination of rocksfrom suitable environments will probably disclose more examples ofnatural nordstrandite.

Nordstrandite from Guam occurs as clusters of tiny crystals radiatingout into microscopic solution cavities in upper Miocene limestone. Cal-cite and red clay occur with the nordstrandite as cavity fillings. Abundantnordstrandite occurs only in the basal part of the limestone near thecontact with old residual soils developed on Eocene and lower Miocenebasalt flows and tuffs. The nordstrandite is younger than the enclosingupper Miocene rocks and probably is much younger.

Hathaway and Schlanger (1965) consider it probable that the'nord-strandite from Guam precipitated from aluminous solutions flowingalong the contact between the limestone and underlying basaltic soil.They suggest that the pH of these solutions did not exceed 8.5 to 9.0because of the limestone environment in which the nordstrandite pre-cipitated. Limestone can be stable, however, at extremely high pHvalues, therefore an upper limit should not be set on the pH of thesesolutions. The enclosing limestone environment can be used, however, toinfer a pH above 7 during nordstrandite crystal growth. ft was in solu-

STRUCTURES OF ALUMINUM HYDROXIDE 7I

tions above pH 7 that nordstrandite formed in our Iaboratory experi-ments thus supporting the inferred conditions of genesis of natural nord-strandite.

Another interesting aspect of the distribution of nordstrandite onGuam is the minor amounts found in limestone some distance from thecontact with volcanic iocks. Hathaway and Carroll (az Schlanger, 1964,p. Daa) speculate that it may be detrital gibbsite that recrystallized tonordstrandite in the limestone environment. The results of our synthesisexperiments support the reasonableness of this hypothesis.

Nordstrandite from west Sarawak, Borneo, occurs as rounded pelletsin terra rossa soil on the edge of a sinkhole in limestone (Wall and others,1962). The authors suggest an origin related to weathering of probableoverlying dacitic sills. Although the data given are meager, we may sur-mise that the limestone host terrain probably limited nordstrandite-precipitating solutions to a pH above 7.

Bayeri,te in Nature. The only well-documented occurrence of naturalbayerite known to us is in Hartruim, fsrael (Bentor and others, 1963;Gross and Heller, 1963). The investigators used X-ray diffraction toidentify the bayerite, which occurs with calcite and gypsum in veinletscutting sedimentary rocks of late Cretaceous age. The rocks are com-posed of calcite and spurrite (a calcium-carbonate-silicate usually formedby contact metamorphism) in a ratio of about two to one. The veinletsalso contain vaterite, portlandite, tobermorite group minerals, thauma-site, and ettringite. Several of these minerals connote genesis at a veryhigh pH (Gross and Heller, 1963), and this agrees with the conditionsunder which pure bayerite f ormed in our experiments (pH l2+).

A recent report of bayerite in a weathering crust of amphibolites andserpentinites in Russia (Khorosheva, 1968) is corroborated by X-raydiffraction. The associated minerals, in this obviously non-equilibriumenvironment, are gibbsite, diaspore, and possibly nordstrandite.

Bayerite-nord.strandite En'igma. The relative ease and speed with whichbayerite, and to a lesser extent nordstrandite, form in the laboratorycontrast with the apparent paucity of these phases in nature. Gibbsite,the slowest forming phase in the laboratory, is by far the most abundantof the three aluminum hydroxide polymorphs in nature.

We have shown that bayerite and nordstrandite form only in alkalinesolutions, and the field evidence implies that the few natural samples ofthese minerals also formed in alkaline environments. Perhaps the scarcityof aluminum-bearing alkaline ground waters is part of the reason for thescarcity of bayerite and nordstrandite in nature. In addition, because

72 ROBERT SCEOEN AND CHARLES E. ROBERSON

of the abundance of silica in most natural alkaline waters, it is temptingto speculate on the importance of silica in favoring the formation ofaluminosilicate minerals rather than aluminum hydroxide.

If silica inhibits the formation of aluminum hydroxide, why does gibb-site form in the pH range from 4 to 7 where silica greatly exceeds alumi-num in most natural waters? The answer may depend upon the dominantform of silica in solution. At pH values below 9, most of the silica is inthe monomeric uncharged form HaSiOr. Above a pH of 9, increasingamounts of silica go into solution as the silicate ion (HsO)Si(OH;ut-(Iler, 1955). Perhaps the charged silicate ion is required to form alumino-silicate minerals, and the virtual absence of this ion at pH values belowabout 7 (except when the silica concentration is high) allows aluminumin solution to precipitate as gibbsite.

If the silicate ion inhibits the formation of bayerite and nordstrandite,the few known natural occurrences of these minerals must representalkaline environments with very little silica in solution. The environmentof nordstrandite deposition on Guam, described by Hathaway andSchlanger (1965), probably possessed a high pH and a low silica concen-tration. Initially, acid ground water in the underlying basalt and tuffshould react with the unstable primary silicates to form secondary sili-cates, increase the pH of the water, and take aluminum into solutionas the aluminate ion (Al(OH)rl-). Little silica should go into solutionbecause these rocks contain almost no uncombined silica (Stark, 1963),and the solubility of the secondary silicates is low in alkaline solutions(Pickering, 1962).

From what we can deduce about the nordstrandite depositional en-vironment in west Sarawak, it resembles that on Guam. Wolfenden(1961) described most of the bauxite deposits in Sarawak as formingfrom intermediate to basic igneous rocks, generally with little quartz.The immediate environment of the nordstrandite pellets on the edge ofa sinkhole in limestone is roughly equivalent to that on Guam.

Barnes and others (1967) reported that ground waters high in pH(lI-12) and low in silica (<6 me/l) form during low temperature ser-pentinization of ultrabasic rocks. Similar reactions in rocks containingaluminous minerals and little free silica may provide ideal conditions forthe precipitation of nordstrandite or bayerite.

We must realize that these nordstrandite occurrences are rare andlimited in extent for good reason. The environment in which nordstrand-ite forms tends to be a small, closed system in which high pH, highaluminum concentration, and low silica concentration can equilibrate.Throughout most of the natural environment, meteoric water in equi-librium with the atmosphere is slightly acid and becomes alkaline only

STRUCTURES OF ALAMINUM EYDROXIDE 73

by reacting with and destroying rock minerals The long-term through-flow of meteoric waters must. result in the destruction of the rock in anenvironment that is slightly acid most of the time. The vastly lowersolubility, in these waters, of aluminum relative to other metals, leadsto the residual enrichment of aluminum hydroxide. And the generallyacid nature of the environment through time causes the gibbsite struc-ture to form.

Our laboratory aging of bayerite-nordstrandite mixtures, althoughunderway for only 2 years, seems to indicate that during aging nord-strandite forms at a greater rate than bayerite (Fig. 8 versus Fig. 6). This,together with the almost complete absence of bayerite in nature, leads usto propose that bayerite is a metastable phase that will eventually recrys-tallize in alkaline solutions to the stable phase nordstrandite. How longit takes for complete recrystallization of bayerite to nordstrandite we donot know, though we detected increased amounts of nordstrandite injust a few months.

If this proposal is correct, the natural bayerite from Hartruim, Israel(Bentor and others, 1963; Gross and Heller, 1963), must be metastable.This natural bayerite formed, presumably, many hundreds or thousandsof years ago. We can speculate that its slow rate of recrystallization tonordstrandite may be due to the extreme dryness of the region in whichit occurs. Hartruim is located on the southwestern edge of the Dead Seain a region with a mean annual rainfall of less than 50 mm (Orni andEftat, t964). 1lhe presence of vaterite (a metastable polymorph ofCaCO) in these bayerite-bearing rocks attests to their dryness, forvaterite is unstable in the presence of water at room temperature andpressure (McConnell, 1960; Johnston and others, 1916).

The uncertainties of extrapolating from conditions of laboratory syn-thesis to natural conditions must always be kept in mind and the pos-sibility of reinterpretation left open. fn the alumina-water system theeffect of most natural foreign ions has been thoroughly studied and foundto slow, but not change, the course of crystallization. Our experiments,and those of others, show that pH is the overriding control on crystalstructure. And, as the preceding discussion shows, the best pH that canbe inferred for the environment of precipitation of natural aluminumhydroxides approximates the pH found in the laboratory.

Suulrany

The mostimportant conclusions deriving from this study are:

1. The pH of the mother liquor controls the structure of the aluminumhydroxide precipitate.

ROBERT SCEOEN AND CHARLES E. ROBERSON

2. The gibbsite structure forms when the pH is less than about 5'8.The bayerite structure forms when the pH is greater than 5.8.

3. The nordstrandite structure forms slowly at pH's greater than5.8. Because nordstrandite seems to form at the expense of previously-

formed bayerite, it may be the stable polymorph of aluminum hydroxidein alkaline solutions.

4. Though the structures of bayerite and nordstrandite are imper-fectly known, they seem to differ from gibbsite and each other solely inthe manner of stacking successive layers of aluminum hydroxide.

5. This difierence in stacking is directly attributable to constraintsimposed by polarization-induced distortions of the hydroxyl ions.

6. The strong and weak distortion of the hydroxyl ions in acid and inalkaline solutions, respectively, is inherited by the gibbsite and bayeritesolids that precipitate from these respective solutions.

7. The structure of nordstrandite seems best explained as a 1:1 regularinterlayering of the gibbsite and bayerite modes of stacking.

8. The suggested stability of nordstrandite relative to bayerite mayreflect a compromise between the higher crystal energy afforded bystrongly polarized hydroxyl ions and the dominance of weakly polarized

hydroxyl ions in the surrounding solution.9. The scarcity of nordstrandite and abundance of gibbsite in nature

reflects, in part, the dominantly acid nature of most weathering environ-ments.

10. The scarcity of nordstrandite, even in alkaline environments, mayindicate a preference by aluminate ions to form aluminosilicates with theusually abundant silicate ions.

The conflict in earlier work regarding the conditions of stability of the

aluminum hydroxide polymorphs arose for two reasons: (1) small quanti-

ties of often poorly crystallized material subject to strong preferred

orientation favored misidentification, and (2) depending upon the pres-

ence of foreign ions and other experimental variables, metastable solidsfrequently formed. The conversion of these metastable phases to stablephases is often slow, especially if the metastable phases are well crys-

Lallized. Metastable products therefore have frequently been consideredto be stable.

We found that syntheses using perchlorate ions, though they bear littlerelation to natural systems, decrease the effect of metastable precursors.

Though the effect is decreased it is not eliminated as shown by our dift-

culty in converting well-crystallized bayerite to nordstrandite duringreasonable periods of aging.

Though a good number of our conclusions represent hypotheses andspeculations based on imprecise data, we feel that these proposals should

STRUCTURES OF ALUMINAM HYDROXIDE 75

be made here so that they may guide future work aimed at their aban-donment, revision, or acceptance.

Acxxowreoournrs

W. L. Polzer and Carol Lind produced the electron micrographs and J. C. Hathawayprovided us with the X-ray difiractogram oI nordstrandite.

We appreciate the efiorts of R. A. Gulbrandsen, J. C. Hathaway, J. D. Hem, and P. H.Hsu, all of whom read the manuscript and suggested changes for its improvement.

We are especially grateful for numerous discussions with J. D. Hem that frequentlysowed the seeds from which our more speculative proposals later blossomed.

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Manuscript ruei.tel,, May 13, 1969; anepteilJor pwbl'ication, f uly 31 , 1969'

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